Hibridization In Methane, Ethane, and Ethylene

BY

Andi Afni Amelia

1213441029

ICP OF CHEMISTRY

MATHEMATIC AND SAINS FACULTYMAKASSAR STATE UNIVERSITY

2012/2013

 Hybridization is the concept of mixing atomic orbitals to form new hybrid

orbitals suitable for the qualitative description of atomic bonding properties. Hybrid

orbitals are very useful in the explanation of the shape of molecular

orbitals for molecules. It is an integral part of valence bond theory. Although

sometimes though together with the valence shell electron-pair repulsion (VSEPR)

theory, valence bond and hybridization are in fact not related to the VSEPR model.

Chemist Linus Pauling first developed the hybridization theory in order to explain the

structure of molecules such as methane (CH4). This concept was developed for such

simple chemical systems, but the approach was later applied more widely, and today

it is considered an effective heuristic for rationalizing the structures of organic

compounds.( http://en.wikipedia.org/wiki/Orbital_hybridisation).

In 1931, linus pauling published the paper in which mathematical models,

utilizing quantum mechanics, explained the shapes of molecules by making use of a

new concept, called hybrid orbitals, in which orbitals of similar energies can “mix”.

For the simplest methane hydrocarbon, formation of bonds requires the uses of the

four sp3 hybrid orbitals formed by combination of one 2s and three 2p orbitals. The

superscript associated with the hybrid designation indicates the number of each

orbital type that has combined to create the hybrid. The number of hybrid orbitals

always equals the number of hybrid orbitals always equals the number of atomic

orbitals that combined.

Because a p orbital has node at the nucleus hybrid orbitals have a contribution

from a p orbital also have a nude at nucleus. We generally ignore the smaller lobe of

the hybrid orbitals because it doesn’t extend as far from the nucleus as the larger one,

and so it’s not involved in the bonding per se. the four sp3 orbital are oriented alones

along lines that pass through the corners of a tetrahedron, as illustrated in figure :

The conceptual combination of the atomic orbitals if carbon that generates sp3

hybrid orbitals. (Sorrel, thomas. 2005, organic chemistry, second edition, university

science books, sausalito, california)

Type of Hibridization

1. Methane

Methane, (CH4, one carbon bonded to four hdrogens) is the simplest organic

molecule. It is a gas at standard temperature and pressure (STP).

This is a flattened, two-dimensional representation of methane that you will

see commonly. The true three-dimensional form of methane does not have any 90

degree angles between bonded hydrogens. The bonds point to the four corners of

a tetrahedron, forming cos-1(-1/3) ≈ 109.5 degree bond angles

(http://en.wikibooks.org/wiki/Organic_Chemistry/Alkanes).

Carbon atomic has two orbitals (2s and 2p) to form ties, meaning if reacts

with hydrogen will be formed two C-H bonds. In fact, the carbon atoms form four

C-H bonds and produce methane molecules with geometrical shapes tetrahedron.

Linus Pauling (1931) describes mathematically how the s orbital and three p orbitals

combine or hybrid atomic orbitals to form four equivalent forms tetrahedral. Orbital

called tetrahedral shaped sp3 hybridization. A three state how many types of orbital

atoms combine, not stating the number of electrons orbitals.

(Stefanus layli prasojo. Kimia organic 1 buku pegangan kimia untuk mahasiswa

farmasi. From http://ashadisasongko.staff.ipb.ac.id/files/2012/02/KIMIA-ORGANIK-

I.pdf). March17th 2013.

Experimentally, methane contains two elements, carbon and hydrogen, and

the molecular formula of methane is CH4.  Both carbon and hydrogen are non-metals,

implying that methane is a covalent compound, not an ionic compound, meaning

methane is made up of molecules, not ions.  According to valence bond theory, the

structure of a covalent species can be depicted using a Lewis structure.  The Lewis

structure of methane is 1, which shows that there are four carbon-hydrogen bonds in

the methane molecule.

(http://science.uvu.edu/ochem/index.php/alphabetical/g-h/hybridization)

Methane

According to valence bond theory, to form a covalent bond, a valence orbital

bearing one electron in one atom overlaps with a valence orbital bearing one electron

in another atom.  Consider the electron configurations of carbon and hydrogen.

The valence shell in carbon is shell two, and it has four electrons. The number

of valence orbitals in carbon each bearing one electron is two (2px and 2py).  The

valence shell in hydrogen is shell one, which has one orbital (1s), bearing one

electron.  Thus, the maximum number of hydrogen atoms a carbon atom can form

covalent bonds with is two, not four, leading to 3, which would be unstable because

the carbon atom lacks an octet of valence electrons.

To explain the structure of the methane molecule, two modifications, known

as excitation and hybridization, are introduced to valence bond theory.

1.  Excitation

The carbon atom used to generate 3 is a ground-state atom.  Convert it into

an excited-state atom by moving an electron from the 2s orbital to the empty

2pz orbital.  Since the latter has higher energy than the former, this change requires

energy.

In the excited-state carbon atom, there are four valence orbitals each bearing

one electron (2s, 2px, 2py, and 2pz).  Thus, the excited-state carbon atom can form

covalent bonds with four hydrogen atoms, resulting in a methane molecule.

2.  Hybridization

Place the four valence orbitals in the excited-state carbon atom used to

generate 4 on a one-dimensional graph in which the single dimension is energy.

Imagine mixing the 2s orbital and the three 2p orbitals to give a homogeneous

mixture and then, dividing the mixture into four new, identical orbitals.  This process

is known as hybridization, and the four new, identical orbitals are called hybridized

orbitals.  Since the four hybridized orbitals are created by mixing one 2s orbital and

three 2p orbitals, they are called sp3-hybridized orbitals, and the energy of an sp3-

hybridized orbital, on the graph, falls between those of a 2s orbital and a 2p orbital,

closer to that of a 2p orbital.  The four atomic orbitals used in hybridization had a

total of four electrons, which are to be distributed in the four sp3-hybridized orbitals.

Since the four sp3-hybridized orbitals have the same energy, according to Hund’s

rule, the electrons must be equally distributed in them.

The shape of the sp3-hybridized orbital is shown mathematically to be roughly as follows. (http://science.uvu.edu/ochem/index.php/alphabetical/g-h/hybridization)

The above diagram is called the orbital diagram. each box in the diagram

orbital states. the relative energies of the various orbitals characterized by the vertical

position of the box in the diagram. electrons are represented by arrows, and the

direction of the spin of the electron is expressed by the direction of the arrow. Four

SP3 hybrid orbitals around the carbon nucleus. because the repulsion between the

electrons in different orbitals, sp3 orbital is located as far as possible from the other

one took extends out of the same carbon nucleus. formed when carbon sp3 bonds, it is

done by overlapping each - one of the sp3 orbitals.

1 orbital sp3 4 orbital sp3 orbital of a hydrogen atom leads to a methane molecule 

(Fessenden, J. R., 1984, Kimia Organik, Edisi Ke-3, Jilid I, Penerbit Erlangga, Jakarta)

2.  Ethane

Two carbons singly bonded to each other with six hydrogens is called ethane.

Ethane (CH3 – CH3) is the second simplest hydrocarbon molecule. It can be

thought of as two methane molecules attached to each other, but with two fewer

hydrogen atoms.

There are several common methods to draw organic molecules. You will use

them interchangeably although sometimes one will work better for one situation or

another (http://en.wikibooks.org/wiki/Organic_Chemistry/Alkanes)

Experimentally, ethane contains two elements, carbon and hydrogen, and the

molecular formula of ethane is C2H6.  Like methane, ethane a covalent compound.

The Lewis structure of ethane is 7, which shows that there are one carbon-carbon

bond and six carbon-hydrogen bonds in the

ethane molecule.

Experimentally, the six carbon-hydrogen bonds in the ethane molecule are identical.

According to VSEPR theory, the geometry at each carbon atom in the ethane

molecule is tetrahedral (8).

There are two similarities between the molecules of methane and ethane.

1. In each molecule, the coordination number of a carbon atom is four.

2. In each molecule, the geometry at a carbon atom is tetrahedral.

The implications is that, to explain the bonding in the ethane molecule using

the valence bond model, two modifications are necessary.

3. Use and excited-state carbon atom, rather than a ground-state carbon atom, in

bonding.

4. Use two sp3-hybridized carbon atoms in bonding.

(http://science.uvu.edu/ochem/index.php/alphabetical/g-h/hybridization)

3. Ethylene

Experimentally, ethylene contains two elements, carbon and hydrogen, and

the molecular formula of ethylene is C2H4.  Like methane and ethane, ethylene is a

covalent compound.  The Lewis structure of ethylene (9) indicates that there are one

carbon-carbon double bond and four carbon-hydrogen bonds in the ethylene

molecule.

Experimentally, the four carbon-hydrogen bonds in the ethylene molecule are

identical.  According to VSEPR theory, the geometry at each carbon atom in the

ethylene molecule is planar trigonal (10).

In the ethylene molecule, each carbon atom forms four bonds, implying that

the ethylene molecule cannot be constructed directly from two ground-state carbon

atoms, for, as explained earlier, a ground-state carbon atom can form a maximum of

only two bonds; only two excited-state carbon atoms can lead to an ethylene

molecule.  The coordination number of each carbon atom in the ethylene molecule is

three, not four as is the case with the carbon atoms in methane and ethane, suggesting

that the two carbon atoms in the ethylene molecule could not possibly be sp3-

hybridized. Imagine mixing the 2s orbital and two of the three 2p orbitals in an

excited-state carbon atom thoroughly to give a homogeneous mixture and, then

dividing the mixture into three new, identical orbitals, which are hybridized,

specifically, sp2-hybridized, orbitals.  Since the three sp2-hybridized orbitals are

created by mixing one 2s orbital and two 2p orbitals, the energy of an sp2-hybridized

orbital falls between those of a 2s orbital and a 2p orbital.  The three atomic orbitals

used in hybridization had a total of three electrons, which are to be distributed in the

three sp2-hybridized orbitals equally, following the Hund’s rule.

The shape of the sp2-hybridized orbital is shown mathematically to be roughly

the same as that of the sp3-hybridized orbital.

To minimize the repulsion between electrons, the three sp2-hybridized orbitals

arrange themselves around the carbon nucleus so that they are as far away from each

other as possible, leading to their planar trigonal arrangement.

The 2p orbital that did not participate in hybridization has one electron.

Again, to minimize the repulsion between electrons, it positions itself so that is as far

away as possible from each sp2-hybridized orbital, leading to the positioning of the 2p

orbital perpendicular to each sp2-hybridized orbital 2(11).  The carbon atom 11 is

called an “sp2-hybridized carbon atom.”

In the ethylene molecule, each carbon atom is bonded to two hydrogen atoms.  Thus, overlap two sp2-hybridized orbitals in 11 with the 1s orbitals of two hydrogen atoms (12).

In the ethylene molecule, each carbon atom is bonded to the other carbon atom, in

addition to two hydrogen atoms.  Thus, take two units of12 and overlap their sp2-

hybridized orbitals, each bearing an electron, to give 13.

13 has the connectivity of the ethylene molecule, but neither carbon atom in 13 has an

octet of valence electrons.  In order to give each carbon atom in 13 an octet of electrons,

overlap the two 2p orbitals laterally to create a pi bond (14).

Each of the four carbon-hydrogen bonds in 14 is formed by the same overlap:

sp2(C)-1s(H).  Consequently, consistent with the observations, the four carbon-

hydrogen bonds in 14 are identical. In addition, in 14, the carbon-carbon bond is a

double bond, one sigma bond and one pi bond, which is consistent with the Lewis

structure of ethylene.  Finally, the geometry at each carbon atom in 14 is planar

trigonal, as predicted by VSEPR theory.  In summary, to explain the bonding in the

ethylene molecule using the valence bond model, two modifications are necessary.

(http://science.uvu.edu/ochem/index.php/alphabetical/g-h/hybridization/).